1. HYDROMETALLURGICAL PROCESS FOR RARE EARTH ELEMENTS RECOVERY FROM
SPENT Ni-HM BATTERIES
Alejandro R. Alonso*(1), Eduardo A. Pérez(2), Gretchen T. Lapidus(2) and Rosa María Luna-Sánchez(1)
(1)Universidad Autónoma Metropolitana Azcapotzalco, Departamento de Energía
Av. San Pablo 180, C.P. 02200México D.F., México
(*Corresponding author: arag@azc.uam.mx)
(2)Universidad Autónoma Metropolitana Iztapalapa, Departamento de Ingeniería de Procesos e
Hidráulica, Av. San Rafael Atlixco 186, C.P. 09340 México D.F., México
ABSTRACT
Rare earth elements have been widely used in various sectors, from the aerospace and steel
industries, to electronic applications, particularly in displays and batteries for portable devices. The high
demand for newer and better equipment, such as cell phones, tablets and lap tops, has increased the
consumption of the rare earth elements and at the same time the need for its recovery from electronic
wastes. For that reason, the recycling of batteries is important, principally the Ni-HM batteries, due to
their elevated content of rare earths elements, in addition to the high concentrations of cobalt and nickel.
The current processes are based on dissolution in strong acid solutions (sulfuric acid concentrations from 2
to 4 mol/L), and temperatures from 30 to 90°C. In the present work, electrodes from spent Ni-HM batteries
were leached using 1 mol/L H2SO4, in the presence of ozone as the oxidant, where recoveries of 96% for
La, Ce and Nd were obtained at room temperature. The separation of Ni and Co from the leach solutions
was performed using an electrochemical reactor, after which rare earth elements were precipitated,
obtaining a mixture of its hydroxides with impurities below 1%, according to EDS analysis.
KEYWORDS
Rare earth elements, leaching, recovery

2. INTRODUCTION
Currently, there is a high demand for energy in portable equipment due to the growing
development of different electronic devices (cellular telephones, digital cameras, PC’s, diverse measuring
equipment, automobiles, etc.). This energy is provided by batteries, which at the end of their useful life,
generate large volumes of waste material. These may contain substances that, even when present in small
quantities, have a harmful effect on the environment; such is the case of chromium, cadmium and lead.
Additionally, considering that practically all of the materials contained in batteries orig inate from non-renewable
sources, the great relevance of the safe handling of this type of waste is evident. In some cases,
recycling these residues may even be profitable.
Batteries are classified as primary (disposable) and secondary (rechargeable). In the case of
secondary batteries, the most important types are based on lithium ion and nickel, which have evolved
since their introduction almost forty years ago in Ni-Fe, Ni-Cd, Ni-HM, Li-ion, ion-polymer. Currently,
these are produced with a large variety of elements; the most important are lithium and nickel compounds,
in addition to cadmium, cobalt, manganese, zinc, iron and some of the rare earth elements, such as cerium,
neodymium, praseodymium and lanthanum.
The rare earth elements (REE) pose a particular problem in that their primary sources are
practically concentrated in one country (Pietrelli et al., 2004), which exports more than 95% of the REE
consumed worldwide. It is also worth noting that they are used in the fabrication of circuits, processors and
memories of a large number of electronic devices. This situation obliges a new approach to materials
recycling of waste batteries, emphasizing the recovery of the elevated REE content, especially in those of
nickel- metal hydride.
Currently, secondary batteries in general are treated by hydrometallurgical techniques to dissolve
the metallic elements that they contain. Zhang et al. (1998) performed studies with hydrochloric acid,
varying the temperature, HCl concentration, solid/liquid ratio and time. With this method, the authors
obtained recoveries of 96% Ni, 100% Co and 99% REE at 95°C with 3 M HCl and a 1:10 solid liquid ratio
in 3 hours. As a less expensive alternative to HCl, Rabah et al. (2007) and Borges et al. (2009) employed
sulfuric acid with hydrogen peroxide as the oxidant; the temperature ranged from 50 to 95°C in > 2 M
H2SO4 using different percentages of peroxide. Recoveries near 90% were achieved for Ni, Co and REE;
however, other authors found that the formation of rare earth sulfates diminished their solubility, even at
95°C (Li et al, 2009). Alternatively, alkaline systems have been tested, taking advantage of the complexing
ability of ammonia (Santos et al., 2012). However, recoveries were lower than those attained in acid
systems.
Some authors have reported the use of complexing agents, such as citrate, in the presence of
sulfuric acid (Innocenzi & Veglio, 2012). Even though this technique improves the Ni and Co solubilities,
favoring REE leaching, it also shifts the reduction potentials for both metals toward more negative values,
incrementing recovery costs.
REE recovery from these liquors should be achieved through precipitation, forming sulfate or
hydroxide salts (Wu et al, 2009) because the reduction potentials are near to -2 V vs SHE (Milazzo, 1978 ),
making their reduction to the metallic state practically impossible in aqueous media. It is worth noting that
these precipitations are not selective because the liquors contain elevated concentrations of Ni and Co
(Santos et al, 2012; Rabah et al, 2007). To avoid co-precipitation, solvent extraction has been employed
with D2EPHA (Yocoyama et al, 1998; Oliveira & Borges, 2009; Tzanetakis & Scott, 2004) and Cyanex
272 (Yocoyama et al, 1998; Oliveira & Borges, 2009; Innocenzi & Veglio, 2012); the latter extractant
separates cobalt from nickel. However, D2EPHA also extracts REE, diminishing the effective separation.
In the present work, a thermodynamic analysis is used to plan the strategy for a three step
separation process. The first used Ni-H batteries that were leached in sulfuric acid solutions employing
ozone as the oxidant. Subsequently, the cobalt and part of the nickel present in the liquor are separated by

3. electrodeposition in a filter press type electrochemical reactor. In the last stage, the REE are precipitated
together with the remainder of the nickel. All processes were performed in sulfuric acid concentrations less
than 1 M.
EXPERIMENTAL METHODOLOGY
Thermodynamic Analysis
Due to the complexity of those systems where many metal ions dissolve, thermodynamic
diagrams were constructed using the MEDUSA software (Puigdomenech, 2010).This software involves an
algorithm that minimizes the free energy of all of the species that form given the declared components
(Ericksson, 1979); the software has its own database, which was enriched with values reported in the
Critical Stability Constants database, NIST 46.8 (2004).
Leaching experiments
The leaching tests were performed in a 500 mL beaker using a solid/liquid ratio of 1:12, mixing at
500 rpm under ambient conditions (T≈25°C and P≈0.78bars). Different leaching systems were studied:
sulfuric, citric and acetic acids in combination with reducing agents, such as hydrazine (N2H4), or oxidants,
such as cupric sulfate (CuSO4) and ozone (O3), in addition to complexing ligands. Table I shows the
composition of the media employed. After leaching, the solids were separated by filtration.
Table I. Leachsolutionscomposition
Leaching/complexing
agent
Concentrations Reducing
agent
Oxidant
agents
0.5M 1M 2M Hydrazine CuSO4 O3
H2SO4 X X X X X
Citric acid X X X
Oxalic acid X X X
Acetic acid X X
EDTA X
EDTA X X
All solutions were prepared with analytical grade reagents and deionized water (11x1018 MΩcm-1).
The reactor shown in Figure 1 was employed when ozone was used as the oxidant. The ozone was
produced by a generator (Basktek, S.A.) and delivered through a porous glass diffusor, maintaining a
constant oxygen flowrate of 1 LPM (336 mg O3/h). The redox potential was controlled with a
potentiometer (Conductronic pH-120) with a combined ORP electrode, to assure that the potential range
was between 0.3-0.8 V vs NHE. After finishing the leach, the solution was filtered to eliminate the
remaining solids.
Dissolution of the metals of interest was monitored at different time intervals and analyzed with
atomic absorption spectrometry (AAS, Varian SpectrAA 220fs) to determine concentration.
Electrochemical tests
The determination of the Co(II) to Co° reduction potential was carried out in a typical three
electrode cell, using a 316 stainless steel disk as the working electrode, with a geometric area of 0.19 cm2.
The auxiliary and reference electrodes were a graphite bar and saturated calomel, respectively. The
potentials reported are all referenced to the standard hydrogen electrode (SHE).

4. Figure 1 -Oxidative leaching system
The cobalt deposit was performed in a parallel plate Electrocell© type reactor. Stainless steel type
316 and titanium plates were used as the cathode and anode respectively, each with an exposed geometric
area of 24 cm2.
The potential determination and the electrodeposition were carried out using a 2263 PARC
potentiostat, connected to a computer with the PowerSuite software for data acquisition and treatment. In
the reactor experiments, samples of the solution were taken from the reservoir, where the changes in
concentration are observed (Alonso, 2007), to be analyzed by AAS.
Precipitation
Selective precipitation was carried out in a 100 mL beaker, in which 50 mL of filtered leaching
solution were placed. Mixing was maintained in ambient conditions. The solution pH was adjusted from
pH 2 to 10 using saturated NaOH and NH4OH solutions. The residence time at each pH value was 1 hour.
The filtered solutions were analyzed before and after each experiment. The solids were rinsed with
100 mL of deionized water and dried at 100°C for 24 hours. Subsequently, they were digested in aqua regia
to determine the elemental composition by AAS.
RESULTS AND DISCUSSION
Thermodynamic Study
Predominance zone diagrams (PDZ) were constructed to determine the conditions at which Ni, Co
and REE were soluble in sulfuric acid solutions. Figures 2 to 4 show the PZD of the Ni(II)SO4
2-,
Co(II)SO4
2- and La(III)SO4
2-systems, respectively. Each metallic ion is shown to form stable complexes
with sulfate in moderate to strong acidic conditions. Ni(II) precipitates at pH values greater than 5.
Figure 2 - Predominance zone diagram (PZD) for the Ni(II)-SO4
2- system

5. On the other hand, Co(II) and La(III) form insoluble hydroxides at pH 8 and 9, respectively.
Special attention should be given to the NiSO4
7H2O precipitate (Figure 2) at sulfate concentrations
greater than 1 M. If the leaching conditions are in this range, the nickel solubilization would probably be
poor, preventing the rupture of the electrode structure; this would probably result in diminished dissolution
also of the other metals present.
Figure 3 - Predominance zone diagram for the Co(II)-SO4
2- system
Figure 4 -Predominance zone diagram for the La(III)-SO4
2- system
The PZD for cerium, praseodymium and neodymium (not shown here) are similar to that of
Figure 4.
Leaching Experiments
The solubility zones predicted in Figures 2 to 4 were used to establish the preliminary leaching
conditions. Considering that nickel is the element found in the highest concentrations in the electrodes, the
leaching conditions were based on the solubility of this metal. For this reason, the sulfuric ac id
concentration was set at 0.5 and 1 M. Table II shows the extraction percentages of each metal considered.
Table II. Recovery of metals as a function of the sulfuric acid concentration
H2SO4
concentration
(mol/L)
Recovery (%)
Ni Co Mn Fe La
0.5 27 42 25 6 >99
1.0 77 >99 >99 22 >99
The total recovery of La and Co, is probably indicative of the structural rupture of the electrode
that contains these elements. However, the relatively low nickel recovery, even at 1 M H2SO4, could be
related to the solubility limit due to the depletion of the reagent by the other metals.

6. The results obtained with 2 M sulfuric acid are presented in Figure 5. Lanthanum recovery was
still 100%, although the cobalt dropped to 93%. Additionally, the maximum nickel dissolution was only
80% after 4 hours, probably due to the formation of solid species at these sulfate concentrations. For that
reason, it would not be convenient to increase the concentration above 1 M H2SO4.
Because nickel is usually found in a reduced state in this type of battery, its recovery should
improve with an oxidative treatment. Tests were performed with 1 M sulfuric acid solutions, adding cupric
sulfate or ozone as oxidant. The elemental dissolution as a function of time with cupric sulfate or ozone is
shown in Figures 6 and 7, respectively. The presence of CuSO4 slightly lowers the extraction of lanthanum,
although that of nickel remains the same as without the oxidant. Because cupric ion cannot be reduced in
this medium (because Cu(I) is not stable), only substitution reaction are possible, which favor the
dissolution of manganese and, to a lesser degree, of cobalt. Furthermore, the increase in the sulfate
concentration favors the formation of insoluble lanthanum compounds, a similar effect to that observed
with higher sulfuric acid concentrations is observed (Figure 5).
100
90
80
70
60
50
40
30
20
10
0
0 1 2 3 4 5
Recovery (%)
Time (hours)
Mn
Ni
Co
Fe
La
Figure 5 - Recovery of metals from Ni-HM spent batteries in a leach using 2 M H2SO4
100
90
80
70
60
50
40
30
20
10
0
0 50 100 150
Recovery (%)
Time (minutes)
Mn
Ni
Co
Fe
La
Figure 6 – Concentration of several elements in a 1 M H2SO4 + 0.2M CuSO4 leaching solution. Conditions:
250 ml solution, 1:12 solid/liquid ratio, 500 rpm, t= 3hrs, T=25°C and P=0.78bars
Figure 7 shows the results obtained for an experiment, where ozone (O3) is continuously sparged
into the acid solution (1 M H2SO4). This produced a solution with a constant redox potential of 700-800
mV. Under these conditions, elevated extractions (> 90%) of Co, La and Mn were achieved almost
immediately. On the other hand, nickel extraction was slower: 80% within the first 60 minutes, gradually
increasing to 100% after 3 hours.
The aforementioned combination of sulfuric acid with ozone resulted in the complete extraction of
REE, Co(II) and Ni(II). However, the procedure for their separation should be carefully pondered. Cobalt
can be electrochemically separated from nickel; however, at the pH of the leach solution (~0.5), hydrogen
evolution is prominent. This factor negatively affects the current efficiency and produces deposits with

7. poor mechanical properties. To minimize hydrogen generation, the pH must be increased, which may
destabilize some of the metal complexes. For that reason, experiments were undertaken to detect any
precipitation of the components of the leaching solution.
100
80
60
40
20
Figure 7 - Metal ions recovery for a 1 M H2SO4 leach with a constant supply of ozone(336 mgO3·h-1)
Conditions: 250 ml solution, 1:12 solid/liquid ratio, 500 rpm, t= 3hrs, T=25°C and P=0.78 bars
Precipitation
The selective precipitation experiments were carried out with solutions from the leach stage.
Concentrated solutions of NaOH or NH4OH were employed to increase the pH. According to published
information, REE are precipitated from leaching solution in the pH range between 0.6 and 2.5 [Li, et al.
2009, Rabah, et al. 2007, Innocenzi, et al. 2012]. It is important to mention that predominance zone
diagram (PZD) shown here for lanthanum (Figure 4), calculated using the NIST database (NIST, 2004),
does not predict the REE precipitation as an oxide.
However, Kim and Osseo-Asare (2012) showed that precipitation of sulfate salts was thermodynamically
possible when large excesses of sulfate are present at pH values above 1. In a preliminary experiment, a
synthetic solution, containing 50% more metal ions than a typical leach liquor (30 g/L Ni, 3.7 g/l Co and
2.8 g/l La [La as the representative element of the REE], was prepared and tested. Figures 8a and 8b shows
the percentage of precipitation of these metals with both NaOH and NH4OH respectively, as a function of
pH. The addition of sodium hydroxide causes significant precipitation of the REE and nickel above pH 2,
which would affect the purity of the solid.
100
80
60
40
20
Figure 8a - Recovery by precipitation as a function of solution pH. Composition: 30 g/l Ni, 3.7 g/l Co and
2.8 g/l La. pH adjusted with: NaOH
0
0 30 60 90 120 150 180
Recovery (%)
Time (hours)
Mn
Ni
Co
Fe
La
0
0 5 10
Precipitation (%)
pH
Ni
Co
La

8. 100
80
60
40
20
0
0 5 10
Precipitation (%)
pH
Ni
Co
La
Figure 8b - Recovery by precipitation as a function of solution pH. Composition: 30 g/l Ni, 3.7 g/l Co and
2.8 g/l La. pH adjusted with: NH4OH
The results from a similar experiment, this time using the leach liquor and adjusting pH with
ammonia (20 g/L Ni, 3.5 g/L Co, 1.4 g/l Mn, 1.7 g/l Fe, 2.65 g/L and 1.75 g/l Nd) are shown in Figure 9.
Nickel, cobalt, manganese and iron present precipitations greater than 30%, but only a t pH values above
pH 6. Both neodymium and lanthanum maintain high percentages of precipitation (> 80%) above pH 1.
The precipitate obtained at pH 2, which represents the value for the best separation of REE, was
rinsed in deionized water and air-dried for XRD analysis. The predominant phases were found to be
cerium, lanthanum, neodymium and praseodymium sulfates (Figure 10). The combined nickel, cobalt and
iron weight was less than 2% for the precipitates formed at pH 2 and 3. In the solids obtained at pH 6 and
7, the nickel content increased to 21 and 97%, respectively. From these results, the recovery of nickel as
Ni(SO4)2∙6H2O at pH 7 seems viable. This only leaves the cobalt to be later recovered by electrodeposition
at pH 7, where the hydrogen evolution is far less.
100
80
60
40
20
0
0 2 4 6
Precipitation(%)
pH
Ni
Co
Mn
Fe
La
Nd
Figure 9 - Recovery by precipitation as a function of solution pH. Leach liquor composition: 20 g/L Ni, 3.5
g/l Co, 1.4 g/l Mn, 1.7 g/L Fe, 2.65 g/l La and1.75 g/l Nd
Electrolysis
Microelectrolysis experiments were performed to determine the best substrate and the potential
windows for nickel and cobalt deposition from the leach liquors. These tests were carried out in a typical
three electrode cell, employing synthetic solutions of 1 M sulfate at pH 7, with and without Co(II) io ns (3.5
g/l), simulating those obtained after the precipitation stage. The substrates examined were 316 SS, high
purity aluminum and commercial titanium. In Figures 11 to 13, typical voltammograms of the solution
with the different electrode materials are shown for 316 SS, Al° and Ti° respectively. The scan was
initiated from the open circuit potential (OCP) in the negative direction. The same scans were performed

9. on solutions without Co (data not shown), observing the reduction of the medium at -500 mV on SS and
Al° and -900 mV on titanium.
La2(SO4)3∙6H2O
Ce2(SO4)3
La2(SO4)3∙6H2O
Ce2(SO4)3
Nd2(SO4)3∙5H2O
Nd2(SO4)3∙8H2
O
Pr2(SO4)3∙8H2O
Ce2(SO4)3∙4H2O
Ce2(SO4)3
Figure 10 - X-Ray diffraction pattern of the precipitate obtained at pH=2
La2(SO4)3∙6H2O
In Figures 11 and 12, a reductive process is observed, commencing at -500 mV, which is probably
due to the reduction of the medium. In contrast, on titanium the formation of reduction wave near -700 mV
(Figure 13) indicates the electrodeposition of cobalt and, probably nickel, followed by the reduction of the
medium at -900 mV. The ability of titanium to shift hydrogen formation to more negative potentials was
taken advantage of in the macroelectrolysis experiments in the parallel plate reactor.
1.E-02
0.E+00
-1.E-02
-2.E-02
-1.1 -0.9 -0.7 -0.5 -0.3 -0.1 0.1
I (mA)
E(V) vs ENH
Figure 11 - Typical voltammograms obtained on 316 SS employing synthetic solutions of 1 M sulfate at
pH 7, with Co(II) ions (3.5 g/l). Scan rate 5 mV/s
The system consisted of an Ecocell type reactor positioned within a continuous recirculation
arrangement. An Ionac 7500 membrane was used to prevent the anodic reactions, such as oxygen
evolution, pH alteration and iron accumulation, from interfering with the cathodic deposition. In this case,
the solution originated from the precipitation stage, with 3.5 g/l cobalt and the remaining 1 g/l nickel. The
applied potential at the cathode was -700 mV, which corresponds to cobalt reduction according to Figure
13. The change in cobalt concentration during the electrolysis is apparently slow (Figure 14), obtaining
only 50% removal in three hours, but with a descending tendency. In contrast, most of the nickel remained

10. in solution (only 10% deposition). This is due to the high stability of nickel in this solution, in addition to
the diminished concentration gradient given that an important part was eliminated in the previous stage.
It should be mentioned that this type of reactor is only useful for validating the energetic
conditions, since it is less accurate for the determination of hydrodynamic factors. This is due principally to
the length/width ratio of the electrodes, in which fully developed velocity profiles are not achiev ed. For
this reason, recoveries typically only reach 50 to 90%. Higher recoveries would be possible in alternative
reactor designs.
5.E-02
0.E+00
-5.E-02
Figure 12 - Typical voltammograms obtained on aluminum, employing synthetic solutions of 1 M
sulfate at pH 7, with Co(II) ions (3.5 g/l). Scan rate 5 mV/s .
8.E-03
3.E-03
-2.E-03
Figure 13 - Typical voltammograms obtained on titanium substrate employing synthetic solutions
of 1 M sulfate at pH. 7, with Co(II) ions (3.5 g/l). Scan rate 5 mV/s
60
50
40
30
20
10
Figure 14 - Cobalt and nickel recovery from the real pregnant solution, during electrolysis in a filter press
type reactor
-1.E-01
-1.1 -0.9 -0.7 -0.5 -0.3 -0.1 0.1
I (mA)
E(V) vs ENH
-7.E-03
-1.1 -0.9 -0.7 -0.5 -0.3 -0.1 0.1 0.3
I (mA)
E(V) vs ENH
0
0 50 100 150
Recovery (%)
Time (minutes)
Ni
Co

11. CONCLUSIONS
The separation of rare earth elements from nickel and cobalt contained in waste secondary
batteries is usually achieved by leaching with high concentrations of sulfuric and/or hydrochloric acids at
temperatures greater than 50°C. However, in the present investigation, it was possible to leach these values
at room temperature, with 1 M sulfuric acid and a constant stream of ozone as the oxidant. This constitutes
the initial stage in a process to recover the dissolved values. Increasing the pH to 3 and 7, permits a
selective precipitation, as sulfate salts, of rare earth elements and nickel, respectively. Once these stages
were concluded cobalt was obtained in its metallic state by electrodeposition using a potential of -700 mV
vs SHE.
REFERENCES
Alonso, A. (2007). Electroseparación selectiva de plata a partir de soluciones amoniacales de tiosulfato.
Tesis de doctorado, UAM-I, D.F., México.
Andricacos, P. C., Arana, C., Tabib, J., Dukovic J & Romankiw, L.T. (1989). Electrodeposition of nickel-iron
alloys. J, Electrochem. Soc., Vol. 136(5) 1136-1340
Berndt, D. (1993). Maintenance- free batteries Lead-Acid, Nickel/Cadmium, Nickel Hydride- A handbook
of battery technology (Electronic & Electrical Engineering Research Studies Power Sources Technology),
Edit. N.E. Bagshaw.
Cabrera, R., González, I., Ávalos J., Vázquez, G.,Pech-Canul, A. (2006). A new approach to describe the
passivity of nickel and titanium oxides. Passivation of Metals and Semiconductors, and Properties of Thin
Oxide Layers, 325-330
Charlot, G. (1967), Chimie analytique générale, Part 1. Paris: Masson & Cie Editeurs.
Cordier, D.J. (2012). Science for a changing world, Rare earths statistics and information. Retrieved from
http://minerals.usgs.gov/minerals/pubs/commodity/rare_earths/.
Eriksson, G. (1979). An algorithm for the computation of aqueous multicomponent, multiphase equilibria,
Anal. Chim. Acta, 112, 375-383.
Fang, W., Sheng-ming, X., Lin-yan, L., Song-zhe, Ch. Gang, X., Jing-ming, X. (2008).Recovery of
valuable metals from anode material of hydrogen-nickel battery. Tran. Nonferrous Met. Soc. China., 19,
468-473.
Gavilán, A., Rojas, L. & Barrera, J. (2009). Las pilas en México: un
diagnóstico ambiental, Instituto Nacional de Ecología. Retrieved from
www2.inecc.gob.mx/publicaciones/consultaPublicacion.html?id_pub=598
Gutiérrez, M. (2012). Lixiviación y recuperación de manganeso a partir de minerales de baja ley. Tesis de
maestría, UAM-I, D.F., México.
Huang, K., Li, J. & Xu, X. (2011), Enhancement of the recycling of waste Ni-Cd and Ni-MH batteries by
mechanical treatment, Waste Management, 31, 1292-1299
Innocenzi, V. & Vegliò, F. (2012). Recovery of rare earths and base metals from spent nickel-metal
hydride batteries by sequential sulphuric acid leaching and selective precipitations. J of Power Sources,
211, 184-191.

12. Ionut, A., Dirk, A. & Hassel, A.W. (2007). Microelectrochemical lithography: A method for direct writing
of surface oxides. Electrochem. Acta, 52, 7865-7869.
Kim, E. & Osseo-Asare, K.W. (2012), Aqueous stability of thorium and rare earth metals in monazite
hydrometallurgy: Eh–pH diagrams for the systems Th–, Ce–, La–, Nd– (PO4)–(SO4)–H2O at 25 °C.
Hydrometallurgy. 113-114, 67-78.
Li, L., Xu, S., Ju, Z. & Wu, F. (2009). Recovery of Ni, Co and rare earths from spent Ni-Metal hydride
batteries and preparation of spherical Ni(OH)2. Hydrometallurgy, 100(1-2), 41-46.
Milazzo, G. (1978) Tables of Standard Electrode Potentials. Ed. John Wiley & Sons, New York, ISBN:
978-0471995340
Oliveira, L.E. & Borges, M. (2009). Hydrometallurgical separation of rare earth elements, cobalt, zinc and
manganese from the positives electrodes of spent Ni-MH batteries from mobile phones. J. of Power
Sources, 195, 3735-3741.
Pietrelli, L., Bellomo,B., Fontan, D. & Montereali M. (2004). Characterization and leaching of NiCd and
NiMH spent batteries for the recovery of metals, Waste Management 25, 221-226.
Pinto J. (2003). Las tierras raras y su relación con los avances tecnológicos de la industria en el presente
siglo. México: Instituto Potosino de Investigación Científica y Tecnológica, 6-C, 1.
Puigdomenech, I. (2010). Make Equilibrium Diagrams Using Sophisticated Algorithms (MEDUSA©),
Royal Institute of Technology, Sweden.
Rabah, M.A., Farghaly, F.E. & Abd-El Motaleb, M.A. (2007).Recovery of nickel, cobalt and some salts
from spent Ni-MH batteries, Waste management, 28, 1159-1167.
Rasmussen, F.E., Ravnkilde, J.T., Tang, P.T., Hansen, O. & Bouwstra, S. (2001). Electroplating and
characterization of cobalt–nickel–iron and nickel–iron for magnetic Microsystems applications. Sensors
and Actuators A., 92, 242-248
Santos, V.E.O., Celante, V.G., Lelis, M.F.F. & Freitas, M.B.J.G. (2012). Chemical and electrochemical
recycling of the nickel, cobalt, zinc and manganese from the positives electrodes of spent Ni-HM batteries
from mobile phones. J. of Power Sources, 218, 435-44
Tzanetakis, N. & Scott, K., (2004). Recycling of nickel–metal hydride batteries. II: Electrochemical
deposition of cobalt and nickel. J. of Chemical Technology & Biotechnology, volume 79(9), 927-934.
Zhang, P., Yokoyama, T., Itabashi, O., Wakui, Y., Suzuki T.M & Inoue, K. (1998). Hydrometallurgical
Process for recovery of metal values from spent nickel metal hydride secondary batteries,
Hydrometallurgy, 50, 61-75.